1 / 70

Chapter Nine

Chapter Nine. Chemical Bonding I : The Lewis Model. Contents. Bonding Models and AIDS Drugs Types of Chemical Bonds Representing Valence Electrons with Dots Ionic Bonding: Lewis Symbols and Lattice Energies Covalent Bonding: Lewis Structures Electronegativity and Bond Polarity

kamal-wiggins
Download Presentation

Chapter Nine

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Chapter Nine Chemical Bonding I : The Lewis Model

  2. Contents • Bonding Models and AIDS Drugs • Types of Chemical Bonds • Representing Valence Electrons with Dots • Ionic Bonding: Lewis Symbols and Lattice Energies • Covalent Bonding: Lewis Structures • Electronegativity and Bond Polarity • Lewis Structures of Molecular Compounds and Polyatomic Ions • Resonance and Formal Charge • Exceptions to the Octet Rule: Odd-Electron Species, Incomplete Octets, and Expanded Octets • Bond Energies and Bond Lengths • Bonding in Metals: The Electron Sea Model

  3. Bonding Models and AIDS Drugs • HIV-protease (protein synthesized by the HIV) is crucial to the virus’s ability to multiply and cause AIDS. • The AIDS drug such as Indinavir, a protease inhibitors, would disable HIV-protease by sticking to the (HIV-protease) molecule’s active site. • Researchers used bonding theories to simulate the shape of potential drug molecules. • Bonding theories • explain (1) how and why atoms attach together to form molecules, (2) why some combinations of atoms are stable and others are not. • predict (1) the shapes of molecules, (2) the chemical and physical properties of compounds

  4. ***** 2. Types of Chemical Bonds • Ionic bond: A chemical bond formed between two oppositely charged ions, generally a metallic cation and a nonmetallic anion, that are attracted to one another by electrostatic forces.

  5. ***** • Covalent bond: A chemical bond in which two atoms share electrons that interact with the nuclei of both atoms, lowering the potential energy of each through electrostatic interactions. • Metallic bond: A type of bond that occurs in metal crystals, in which metal atoms donate their electrons to an electron sea, delocalized over the entire crystal lattice.

  6. 3. Representing Valence Electrons with Dots • Lewis dot symbols: Combination of a chemical symbol for the element and valence electrons (the dots around the element symbol) of the atom. • The oxygen element O for example:

  7. ***** • Lewis dot symbols for representative elements: • The first four dots are placed singly on each of the four sides of the chemical symbol, and pairing the dots as the next four are added.

  8. Lewis Bonding Theory • Lewis structures: A drawing that represents chemical bonds between atoms as shared or transferred electrons. • Metals and Nonmetals combine: valence electrons transferred from the metal to the non-metal atoms giving rise to ionic bonds. • Nonmetals and Nonmetals combine: one or more pairs of valence electrons are shared between the bonded atoms producing covalent bonds.

  9. Continued • Octet Rule: When atoms bond, they tend to gain, lose, or share electrons to result in eight valence electrons (ns2np6, noble gas configuration). • Exceptions: • H, Li, Be, B attain an electron configuration like He (a duet). • Expanded octets for elements in period 3 or below, using empty valence d orbitals.

  10. 4. Ionic Bonding: Lewis Symbols and Lattice Energies • Lewis structure for ionic bonding KCl for example: more example: Na2S ***** more example: CaCl2

  11. Example 9.1Using Lewis Symbols to Predict the Chemical Formula of an Ionic Compound Use Lewis symbols to predict the formula for the compound that forms between calcium and chlorine. Solution Draw Lewis symbols for calcium and chlorine based on their number of valence electrons, obtained from their group number in the periodic table. Calcium must lose its two valence electrons (to be left with an octet in its previous principal shell), while chlorine only needs to gain one electron to get an octet. Draw two chlorine anions, each with an octet and a 1– charge, and one calcium cation with a 2+ charge. Place brackets around the chlorine anions and indicate the charges on each ion. Finally, write the formula with subscripts to indicate the number of atoms. CaCl2

  12. Lattice Energy • Lattice energy: The energy associated with forming a crystalline lattice from gaseous ions. Lattice energy of NaCl for example: Na+(g)+ Cl–(g) → NaCl(s)ΔH = Lattice energy = –788 kJ The large negative value of the lattice energy (exothermic) makes ionic bonding.

  13. Continued • Born-Haber cycle: A hypothetical series of steps based on Hess’s law that represents the formation of an ionic compound from its constituent elements. Formation of NaCl from elements Na and Cl for example:

  14. ***** Continued • Calculating lattice energy from Born-Haber cycle (NaCl for example): 1. Enthalpy of sublimation: Na(s) → Na(g)ΔH1 = +108 kJ 2. Bond energy: ½Cl2(g) → Cl(g)ΔH2 = +122 kJ 3. First ionization energy: Na(g) → Na+(g) + e–ΔH3 = +496 kJ 4. Electron affinity: Cl(g) + e– → Cl–(g)ΔH4 = – 349 kJ 5. Lattice energy: Na+(g) + Cl–(g) → NaCl(s) Lattice energy = ΔH5 = ? Overall Enthalpy of standard formation: Na(s) + ½ Cl2(g) → NaCl(s)ΔHfo = –411 kJ ΔHfo = ΔH1 + ΔH2 + ΔH3 + ΔH4 + ΔH5 = –411 kJ Lattice energy = ΔH5 = ΔHfo – (ΔH1 + ΔH2 + ΔH3 + ΔH4) = –788 kJ

  15. Trends in Lattice Energies: Ion Size Alkali metal chlorides for example: * Lattice energies become less exothermic (less negative) with increasing ionic radius.

  16. Trends in Lattice Energies: Ion Charge NaF and CaO for example: * Lattice energies become more exothermic (more negative) with increasing magnitude of ionic charge.

  17. ***** Example 9.2Predicting Relative Lattice Energies Arrange these ionic compounds in order of increasing magnitude of lattice energy: CaO, KBr, KCl, SrO. Solution KBr and KCl should have lattice energies of smaller magnitude than CaO and SrO because of their lower ionic charges (1+, 1– compared to 2+, 2–.) When you compare KBr and KCl, you expect KBr to have a lattice energy of lower magnitude due to the larger ionic radius of the bromide ion relative to the chloride ion. Between CaO and SrO, you expect SrO to have a lattice energy of lower magnitude due to the larger ionic radius of the strontium ion relative to the calcium ion. Order of increasing magnitude of lattice energy: KBr < KCl < SrO < CaO

  18. Ionic Bonding: Models and Reality • All ionic compounds are solids at room temperature. • Ionic compounds have high melting points and boiling points, MP generally > 300 °C. • Ionic solids are relatively hard (compared to most molecular solids) • Ionic solids are brittle. When struck they shatter. • Ionic solids do not conduct electricity. • Ionic compounds conduct electricity in the liquid state or when dissolved in water (aqueous state).

  19. 5. Covalent Bonding: Lewis Structures • Lewis Structure for Simple Molecules • Lewis structure of molecule shows the bonded atoms with the electron configuration of a noble gas. • The shared electrons is counted for each atom that shares them. • The valence electrons of the bonded atoms obey the octet rule (H obeys the duet rule). • bonding pairs: The shared pairs of electrons (two electrons) in a molecule, commonly represented by a dash (—). • Lone pairs (nonbonding pairs): The electron pairs which are not shared.

  20. Single Covalent Bonds When two atoms share one pair of electrons, it is called a single covalent bond (two electrons), represented by (–). Example: H2O ***** Example: Cl2 Example: H2 *****

  21. Double Covalent Bonds When two atoms share two pairs of electrons the result is called a double covalent bond (four electrons), represented by (=). Example: O2 ***** • Triple Covalent Bonds When two atoms share three pairs of electrons the result is called a triple covalent bond (six electrons), represented by (). Example: N2 *****

  22. ***** • Non-metals of the second period (except boron) form a number of covalent bonds equal to eight minus the group number:

  23. Covalent Bonding: Models and Reality • The covalent bonds are highly directional and the fundamental units of covalently bonded compounds are individual molecules. • Covalently bonded molecular compounds: the interactions between molecules (intermolecular forces) are generally much weaker than the bonding interactions within a molecule (intramolecular forces).

  24. Intermolecular and Intramolecular Forces * To boil a molecular substance, one simply have to overcome the relatively weak intermolecular forces, so molecular compounds generally have low boiling points.

  25. Continued • Molecular compounds are found in all three states at room temperature. • Molecular compounds have low melting points and boiling points, MP generally < 300 °C. • Molecular compounds do not conduct electricity in the solid or liquid state. • Molecular acids conduct electricity when dissolved in water (aqueous state) due to them being ionized by the water.

  26. ***** • Electronegativity and Bond Polarity • Electronegativity (EN): The ability of an atom to attract bonding electrons to itself. • EN is derived from ionization energy and electron affinity. • The greater the EN of an atom, the more strongly it attracts the electronsin a chemical bond. • EN generally increases across a period in the periodic table. • EN generally decreases down a column in the periodic table.

  27. ***** • Electronegativity Difference on Bond Type • Electronegativity Difference (ΔEN) : The difference in EN of bonded atom • ΔEN versus Bond type: • Pure nonpolar covalent bond: ΔEN = 0, ex. Cl─Cl. • metal-metal (metalic) bond, all atoms with low EN

  28. ***** • Ways for expressing electron distributions of covalent bond δ+ δ– H—H H—Cl H—Cl • Polyatomic ion such as CO32–: Covalent bond between C and O • Ionic compounds such as Na2CO3: Ionic bond between Na+ and CO32–

  29. Polar molecule in electric field Oriented randomly Oriented regularly

  30. Dipole Moment (μ) • A measure of the separation of positive and negative charge in a molecule. • A diatomic molecule with a polar covalent bond (such as HCl) for example: μ = δd δ: magnitude of the partial charge (unit: coulomb, C) d: distance that separates positive and negative charge (unit: m) Unit for μ: Debye (D), 1 D = 3.34 x 10–30 C m

  31. Dipole Moment versus ΔEN

  32. Percent Ionic Character • The percentage of a bond’s measured dipole moment compared to what it would be if the electrons were completely transferred, a reference μ with 6.2 D.

  33. Percent Ionic Character versus ΔEN

  34. ***** • Bond Type, ΔEN, and Percent Ionic Character A Summary

  35. Example 9.3Classifying Bonds as Pure Covalent, Polar Covalent, or Ionic Determine whether the bond formed between each pair of atoms is covalent, polar covalent, or ionic. a. Sr and F b. N and Cl c. N and O Solution a. In Figure 9.8, find the electronegativity of Sr (1.0) and of F (4.0). The electronegativity difference (ΔEN) is ΔEN = 4.0 – 1.0 = 3.0. Using Table 9.1, classify this bond as ionic. b. In Figure 9.8, find the electronegativity of N (3.0) and of Cl (3.0). The electronegativity difference (ΔEN) is ΔEN = 3.0 – 3.0 = 0. Using Table 9.1, classify this bond as covalent. c. In Figure 9.8, find the electronegativity of N (3.0) and of O (3.5). The electronegativity difference (ΔEN) is ΔEN = 3.5 – 3.0 = 0.5. Using Table 9.1, classify this bond as polar covalent.

  36. ***** • Lewis Structures of Molecular Compounds and Polyatomic Ions • Steps: • Write a skeletal structure and connect the atoms by single dashes (covalent bonds). • Determine the total number of valence electrons. • Place lone pairs electrons around the terminal atoms to give each terminal atom (except H) an octet. Then, place the remaining electrons around the central atom to give an octet. • If necessary, move one or more lone pairs of electrons from terminal atom to form a multiple bond to the central atom.

  37. ***** • Skeletal structure (the arrangement of atoms): • Hydrogen atoms are usually terminal atoms • Either molecules or polyatomic ions usually have compact and symmetrical structures. • The central atom of a structure usually has the lowest electronegativity. • Oxoacids such as HClO4, HNO3, etc., hydrogen atoms are usually bonded to oxygen atoms

  38. Example 9.4Writing Lewis Structures Write the Lewis Structure for CO2. Procedure For… Writing Lewis Structures for Covalent Compounds Solution Step 1 Write the correct skeletal structure for the molecule. Because carbon is the less electronegative atom, put it in the central position. O C O Step 2 Calculate the total number of electrons for the Lewis structure by summing the valence electrons of each atom in the molecule. Step 3 Distribute the electrons among the atoms, giving octets (or duets for hydrogen) to as many atoms as possible. Begin with the bonding electrons, and then proceed to lone pairs on terminal atoms, and finally to lone pairs on the central atom. Bonding electrons are first. O:C:O (4 of 16 electrons used)

  39. Example 9.4Writing Lewis Structures Continued Lone pairs on terminal atoms are next. (16 of 16 electrons used) Step 4 If any atom lacks an octet, form double or triple bonds as necessary to give them octets. Since carbon lacks an octet, move lone pairs from the oxygen atoms to bonding regions to form double bonds.

  40. Example 9.5Writing Lewis Structures Write the Lewis Structure for NH3. Procedure For… Writing Lewis Structures for Covalent Compounds Solution Step 1 Write the correct skeletal structure for the molecule. Since hydrogen is always terminal, put nitrogen in the central position. Step 2 Calculate the total number of electrons for the Lewis structure by summing the valence electrons of each atom in the molecule.

  41. Example 9.5Writing Lewis Structures Continued Step 3 Distribute the electrons among the atoms, giving octets (or duets for hydrogen) to as many atoms as possible. Begin with the bonding electrons, and then proceed to lone pairs on terminal atoms, and finally to lone pairs on the central atom. Bonding electrons are first. (6 of 8 electrons used) Lone pairs on terminal atoms are next, but none is needed on hydrogen. Lone pairs on central atom are last. (8 of 8 electrons used) Step 4 Since all of the atoms have octets (or duets for hydrogen), the Lewis structure for NH3 is complete as shown in the previous step.

  42. Example 9.6Writing Lewis Structures for Polyatomic Ions Write the Lewis structure for the NH4+ ion. Solution Begin by writing the skeletal structure. Since hydrogen is always terminal, put the nitrogen atom in the central position. Calculate the total number of electrons for the Lewis structure by summing the number of valence electrons for each atom and subtracting 1 for the 1+ charge. Place two bonding electrons between every two atoms. Since all of the atoms have complete octets, no double bonds are necessary. (8 of 8 electrons used)

  43. Example 9.6Writing Lewis Structures for Polyatomic Ions Continued Lastly, write the Lewis structure in brackets with the charge of the ion in the upper right-hand corner.

  44. ***** 8. Resonance and Formal Charge • Resonance • Resonance structures: A molecule or a polyatomic ion is presented two or more valid Lewis structures that are shown with double-headed arrows between them. • Resonance hybrid: The actual structure of a molecule or polyatomic ion that is intermediate between two or more resonance structures. • O3 for example:

  45. ***** Continued • Electrons in resonance hybrid are spread out over several atoms, called delocalized electrons. • The resulting stabilization of the electrons (that is, the lowering of their potential energy due to delocalization) is sometimes called resonance stabilization.

  46. Example 9.7Writing Resonance Structures Write a Lewis structure for the NO3– ion. Include resonance structures. Solution Begin by writing the skeletal structure. Since nitrogen is the least electronegative atom, put it in the central position. Calculate the total number of electrons for the Lewis structure by summing the number of valence electrons for each atom and adding 1 for the 1– charge. Place two bonding electrons between each pair of atoms. (6 of 24 electrons used)

  47. Example 9.7Writing Resonance Structures Continued Distribute the remaining electrons, first to terminal atoms. There are not enough electrons to complete the octet on the central atom. (24 of 24 electrons used) Form a double bond by moving a lone pair from one of the oxygen atoms into the bonding region with nitrogen. Enclose the structure in brackets and include the charge.

  48. Example 9.7Writing Resonance Structures Continued Since the double bond can form equally well with any of the three oxygen atoms, write all three structures as resonance structures. (The actual space filling model of NO3– is shown here for comparison. Note that all three bonds are equal in length.)

  49. ***** • Formal Charge • Formal Charge: A fictitious charge that an atom in a Lewis structure would have if all the bonding electrons were shared equally between the bonded atoms. • Estimation expression: Number ofvalence electronsin the unbonded atom Number ofbonding electrons on the bonded atom Number oflone-pair electronson the bonded atom – – ½ Simply expressed by: FC = V – LP – ½ BP • The sum of all formal charges in a neutral molecule must be zero, in a polyatomic ion must equal the charge of the ion.

More Related